1 / 16

Lewis Dot Structures for Bonding

Lewis Dot Structures for Bonding. Review. What is bonding? Why does it occur? How do they form? (Explain difference between ionic and covalent. What is bonding? Attraction of valence electrons of one atom to the nucleus of another atom Why does it occur?

denton-buck
Download Presentation

Lewis Dot Structures for Bonding

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Lewis Dot Structures for Bonding

  2. Review • What is bonding? • Why does it occur? • How do they form? (Explain difference between ionic and covalent

  3. What is bonding? • Attraction of valence electrons of one atom to the nucleus of another atom • Why does it occur? • STABILITY!!!! Because of bonding, each element will obtain full valence/outer shell (8 e-) • How do they occur • Ionic (metal and non metal) • Metal gives electrons to nonmetal • Atoms are then attracted, but DO NOT OVERLAP

  4. Covalent Bonding Two non metals SHARE a pair of electrons Called a molecule Shared electrons will allow both atoms to obtain the octet in valence shell Electrons will be shared equally between the two atoms

  5. How do they form? • Overlap of orbitals is how the electrons are shared • The overlap causes a pseudo octet in valence • Both atoms feel like they have an octet • Two types of overlap • Sigma • pi

  6. Sigma bond • Overlap of two s orbitals • Share only one pair of electrons • Called a single bond

  7. Pi bond • Overlap of p orbitals • Sharing of two or three pairs of electrons • Share two pairs of electrons = double bond • Share three pairs of electrons = triple bond

  8. Which bonds are stronger?? Shorter?? • The more electrons that are shared the strongerthe bond will be. • Bond energy – energy required to break the bond • The more electrons that are shared the tighter/shorterthe bond will be • Single • Double • Triple Weakest, longest Strongest, shortest

  9. How can we represent these bonds? • Lewis dot structures • How do we write these for individual atoms? • N? Cl? • Number of non-paired e- = number of bond possible F F • Circled dots represent the shared electron pair • Other electron pairs are not involved in the bonding • Belong to one atom – nonbonding pair

  10. How do I do this??? • Step 1:Determine types of atom and how many you will need – i.e. the formula • PH3 • Step 2:Write Lewis dot notations for individual atoms • P? H?

  11. Step 3:Determine total number of valence electrons in compound • Add together the number of valence electrons for each element P: 1 x 5e- = 5e- H: 3 x 1e- = 3e- 8e-

  12. Step 4: Least electronegative element goes in center or element with most bonding sites Except: - Hydrogen NEVER in center - If carbon is present, it goes in the center H H P H • Step 5: Place an electron pair between every element H H P H

  13. Step 6: Add unshared pairs to all atoms until they have a full Octet (hydrogen should only have 2) H H P H • Step 7: Self check, number of electrons should equal match total number of valence electrons.

  14. Let’s do another example • NI3

  15. Lewis Dot for double and triples bonds • Follow steps 1 – 7 • If there are too many electrons, the molecule may have double or even triple bonds. • Step 8:Remove electron pairs until desired number of electrons is reached. • Step 9:Some elements will no longer have an octet so we must move e- pairs to non-hydrogen bonds until all elements have an octet

  16. Examples • CO2 • HCN

More Related